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Paper chromatography: overview

Paper chromatography is an analytical technique used for separating and identifying components of a mixture. The most common form of paper chromatography is ascending paper chromatography, which involves applying a small drop of a solution mixture to a piece of filter paper a short distance from one end. After drying the drop of solution, the paper is suspended in a container, with the same end of the paper dipped in a solvent without immersing the drop itself.

The mixture is separated as the solvent flows up the paper by capillary action, moving the components of the mixture at different rates. The diagram below shows a drop of black ink being separated into its coloured components.

Examples of other forms of paper chromatography are descending chromatography and radial chromatography.

 

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Acids and bases: overview

The word acid comes from acidus, a latin word meaning sour while the term ‘base’ was introduced by Guillaume Rouelle, a French chemist, in 1754. Since then, scientists have developed several concepts of acids and bases and some related definitions. These concepts and definitions serve to help us understand and predict reactions involving acids and bases.

 

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Definitions of acids and bases

What are the different definitions of acids and bases?

Svante Arrhenius (in 1884)

Arrhenius acid:    A substance, when in its aqueous state, produces hydrogen ions, H+.

Arrhenius base:  A substance, when in its aqueous state, produces hydroxide ions, OH–.

For example, aqueous hydrogen chloride (or hydrochloric acid) is an acid:

HCl(aq)\rightarrow H^{+}(aq)+Cl^{-}(aq)

However, these definitions are restrictive, for example, they do not consider gaseous hydrogen chloride and gaseous ammonia as an acid and a base respectively.

NH_3(g)+HCl(g)\rightarrow NH_4Cl(s)

 

Bronsted-Lowry (in 1923)

Bronsted acid:     A substance that is a H+ (proton) donor.

Bronsted base:    A substance that is a H+ acceptor.

Bronsted-Lowry’s definitions expand on Arrhenius’. Gaseous hydrogen chloride and gaseous ammonia are now an acid and a base respectively.

HCl(g)+H_2O(l)\rightarrow H_3O^+(aq)+Cl^-(aq)

NH_3(g)+H_2O(l)\rightarrow NH_4^+(aq)+OH^-(aq)

Gaseous hydrogen chloride donates a proton to water to form the hydronium ion, H3O+ while gaseous ammonia accepts a proton from water to form the ammonium ion, NH4+. Note that the hydrogen ion is more likely to exist as a hydronium ion (or larger complexes like H9O4+) in water than as an isolated proton, H+. The hydronium and ammonium ions formed are acids too as they can donate their protons to bases. Bronsted and Lowry called them conjugate acids. Similarly, the chloride and hydroxide ions are bases as they can accept protons from acids. They are called conjugate bases. In general,

\begin{matrix} HA\\acid \end{matrix}\: \: \: \: +\: \: \: \:\begin{matrix} B\\base \end{matrix}\: \: \: \: \rightarrow \: \: \: \:\begin{matrix} BH^+\\conjugate\: acid \end{matrix}\: \: \: \: +\: \: \: \:\begin{matrix} A^-\\conjugate\: base \end{matrix}

The HA/A– and BH+/B pairs are called conjugate acid-base pairs.

 

Question

Is OH– in the following reaction a base or an acid?

O^{2-}(aq)+H_2O(l)\rightarrow 2OH^-(aq)

Answer

O2- is a base and H2O is an acid. OH– is a conjugate acid of O2- and a conjugate base of H2O.

 

Acids with one, two and three dissociable hydrogen atoms are known as monoprotic (e.g. HCl), diprotic (e.g. H2SO4) and triprotic (e.g. H3PO4) respectively. Acids with more than one dissociable hydrogen atoms are collective called polyprotic acids. Similarly, bases that can accept one, two and many protons are called monoprotic bases, diprotic bases and polyprotic bases respectively. Finally, some substances can be both proton acceptors and proton donors, for example, the hydrogen carbonate ion, HCO3– :

HCO_3^-(aq)+H_3O^+(aq)\rightarrow H_2CO_3(aq)+H_2O(l)

HCO_3^{\: -}(aq)+OH^-(aq)\rightarrow CO_3^{\: 2-}(aq)+H_2O(l)

Such substances are called amphiprotic.

 

Gilbert Lewis (in 1938)

Lewis acid:          A substance that is an electron pair acceptor.

Lewis base:        A substance that is an electron pair donor.

According to Lewis’ definitions, HCl is an acid and NH3 is a base because HCl accepts a pair of electrons from the nitrogen atom of NH3 (the electron pair donor) to give the ammonium and chloride ions.

Lewis’ definitions not only encompass substances under Bronsted-Lowry’s definitions but also include species that do no H, e.g. BF3.

 

Question

What is the difference between a base and an alkali?

Answer

An alkali is a hydroxide of a group 1 (alkali metal) or group 2 (alkaline earth metal) element with pH > 7 in its aqueous form, e.g. NaOH and Mg(OH)2. A base is defined either by an Arrhenius base, a Bronsted base or a Lewis base. In other words, an alkali is a soluble hydroxide, while a base can be any compound that fits the definition of either an Arrhenius base, a Bronsted base or a Lewis base, e.g. ammonia gas, NaH and any soluble hydroxide. Therefore, we can also define an alkali as the solution of a water-soluble base.

\begin{matrix} NH_3(g)\\base \end{matrix}+H_2O(l)\rightarrow \begin{matrix} NH_4OH(aq)\\alkali \end{matrix}

\begin{matrix} NaH(s)\\base \end{matrix}+H_2O(l)\rightarrow \begin{matrix} NaOH(aq)\\alkali \end{matrix}+H_2(g)

 

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The pH scale

The pH scale is a range of numbers for measuring the extent of acidity or basicity of a substance.

As mentioned in the previous article, pH is defined as –log[H+] and is accurate over the range of 0 to 14, which is equivalent to concentrations of aqueous H+ of between 1 M to 10-14 M. We can use the values of 0 to 14 to construct a colour-coded pH scale as a convenient reference of the concentration of H+ in solutions.

So, instead of saying that the concentration of H+ in a sample of urine is 10-6 mol dm-3, we simply say that the sample of urine has a pH of 6.

Question

What if the concentration of H+ in a solution is 2 M? Does it mean that the pH of the solution is -0.3?

Answer

We often use a pH glass electrode to measure the pH of an aqueous solution. Although it is not wrong to say that the pH of a 2 M HCl solution is -0.3, the pH scale becomes less accurate when the concentration of a solution is greater than 1 M. This is because acid molecules may be absorbed by a layer of gel, which is coated on the glass bulb (see below diagram), lowering the activity of H+ at the electrode, making it harder to accurately measure the real concentration of H+ in the solution. On the other hand, a solution with pH < 10-14 M has so low a concentration of H+ that other ions present in the solution may be detected by the electrode instead, giving an inaccurate reading. This is why the pH scale is usually quoted between 0 to 14. Fortunately, many common solutions like those stated in the diagram above have pH values within the measurable range.

 

 

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Definition of pH

pH, which stands for power of hydrogen, is a measure of the acidity of an aqueous solution. It is defined as:

pH=-log[H^+]\; \; \; \; \; \; \;\; (1)

where [H+] is the concentration of the aqueous hydrogen ion.

The concept was developed by the Danish chemist, Soren Peder Lauritz Sorensen in 1909.

Question

Why is pH defined on a logarithmic scale and why does it have a negative sign in front?

Answer

The most accurate way of measuring the concentration of H+ in a solution is using a glass electrode, a type of ion-selective electrode (see below diagram). The voltage, E, of the electrochemical cell containing the electrode is measured at 25oC and the concentration of H+ is determined by the Nernst equation:

E=0.059log[H^+]\; \; \; \; \; \; \; \; (2)

Scientists found that the values generated by the factor log[H+] at different voltages are accurate over the range of -14 to 0 and decided to use it to conveniently express the concentration of H+ in a solution. As it is easier to deal with positive numbers (i.e. 0 to 14 instead of -14 to 0), pH is defined as –log[H+], which makes eq2:

E=-0.059pH

 

 

 

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Acidity vs alkalinity

The balance between acidity and alkalinity plays a crucial role in everything from our health to the environment, influencing processes as varied as digestion and climate stability.

Pure water ionises to give hydrogen and hydroxyl ions with the concentrations of H+ and OH– at 10-7 M each.

H_2O\rightleftharpoons H^++OH^-

This makes the pH of pure water 7. Water also ionises in aqueous solutions (i.e. solutions of compounds, e.g. NaCl, where water is the solvent). We call an aqueous solution of pH 7, a neutral solution. If the pH of an aqueous solution is less than 7, [H+] > [OH–] and the solution is termed acidic. If the pH of the aqueous solution is more than 7, [H+] < [OH–] and the solution is basic or alkaline.

 

Question

20.0 ml of 0.0025 M of HCl is added to 50.0 ml of water. What is the pH of the final solution? Why is it not necessary to consider H+ from the dissociation of water in the computation?

Answer

No. of moles of H+ in 20 ml of HCl = \frac{0.0025\times 20}{1000}

[H+] in 70 ml of solution = \frac{0.0025\times 20}{1000}/0.070

pH=-log(\frac{0.0025\times 20}{1000}/0.070)=3.15

We do not consider the concentration of H+ from water as it is insignificant. According to Le Chateliers’ principle, the presence of H+ from aqueous HCl causes the equilibrium of the dissociation of water to shift to the left, reducing the concentration of H+ in water to less than 10-7 M. Even if we assume that , the resulting pH value of -log\left \{[(\frac{0.0025\times 20}{1000})+(10^{-7}\times \frac{50}{1000})]/0.070\right \} is still 3.15.

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pOH, pKw, pKa and pKb

Understanding pOH, pKw, pKa, and pKb is essential for grasping the intricacies of acid-base chemistry, as they define the relationships between concentrations of hydrogen ions and their corresponding bases in aqueous solutions.

Extending on the concept of pH, we can define the negative logarithm of [OH–], Kw, Ka and Kb respectively as follows:

pOH=-log[OH^-]\; \; \; \; \; \; \; \; (3)

pK_w=-logK_w\; \; \; \; \; \; \; \; (4)

pK_a=-logK_a\; \; \; \; \; \; \; \; (5)

pK_b=-logK_b\; \; \; \; \; \; \; \; (6)

Since Kw = [H+][OH–],

-logK_w=-log[H^+]-log[OH^-]\; \; \; \; \; \; \; \; (7)

Substitute eq1, eq3 and eq4 in eq7

pK_w=pH+pOH

At 250C, Kw = 10-14. So, pKw = 14 and

pH+pOH=14

 

Question

The concentration of hydroxide ions in the urine of a patient with kidney stones is 1.698 x 10-10 M. What is the pH of the patient’s urine?

Answer

pOH=-log(1.698\times 10^{-10})

pH=14-[-log(1.698\times 10^{-10})]=4.23

 

 

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Titration: overview

Titration, or titrimetry, is an analytical technique for determining the concentration of a chemical substance called an analyte by reacting it with a solution of known concentration.

There are many different titration methods, including acid-base titration, redox titration, complexometric titration, gravimetry titration, potentiometric titration, coulometric titration, thermometric titration and radiometric titration.

This article focuses on acid-base titration and redox titration. In an acid-base titration (e.g., HCl and NaOH), if the analyte is an acid, the reacting solution is the base, which is added in consecutively small amounts (using the apparatus shown in the diagram below) to the other. The analyte can be either in the burette (the dispensing device) or in the conical flask.

During the preparation stage of the titration, the burette is filled with an acid (or base) through a glass funnel, which is inserted into the burette’s top opening. A pipette (a chemical dropper) is used to introduce the base (or acid) to the conical flask. A few drops of an auxillary reagent called an indicator (a colour changing compound) is also added before start of the reaction. The reaction begins when small amounts of the reagent is added from the burette to the substance in the conical flask and is complete when a significant change of colour of the solution in the flask is observed. The concentration of the analyte is then calculated via stoichiometric analysis by noting the volume of solution added from the burette to the flask.

For example, a solution in a flask containing 25.0 cm3 of an unknown concentration of NaOH with a few drops of methyl orange (indicator) appears red at the beginning of the titration. The solution remains red as 0.1 M HCl is gradually added from the burette. Just before reaching the equivalence point, the solution turns orange, which signals to the person conducting the experiment that the next couple of drops of HCl will turn the solution yellow (the end point), where the titration is complete. The concentration of NaOH is calculated with the volume of HCl used (e.g. 12.5 cm3) and by referring to the neutralisation equation:

HCl(aq)+NaOH(aq)\rightarrow NaCl(aq)+H_2O(l)

According to the stoichiometry of the above reaction, 0.1\times \frac{12.5}{1000} moles of HCl react with 0.1\times \frac{12.5}{1000} moles of NaOH. Therefore,

[NaOH]=\left (0.1\times \frac{12.5}{1000}\right )/\frac{25.0}{1000}=0.05\, M

In a redox titration (e.g. Fe2+ and MnO4– with phosphoric (V) acid), the same procedures apply but the indicator may not be required as either the analyte or the solution of known concentration may have the ability to change colour to pinpoint the completion of the reaction.

\begin{matrix} MnO_4^{\; -}(aq)\\purple \end{matrix}+5Fe^{2+}(aq)+8H^+(aq)\rightarrow \begin{matrix} Mn^{2+}(aq)\\colourless \end{matrix}+5Fe^{3+}(aq)+4H_2O(l)

In the above titration, MnO4– is added from the burette to Fe2+. Therefore, the solution in the flask remains colourless until the end point, where the solution turns pale pink.

Question

Why is an acid added to a permanganate-to-iron titration and why is H3PO4 the preferred acid as compared to HCl?

Answer

The oxidising strength of MnO4– varies with the pH of its solution. At low and high pH, MnO4– is reduced by Fe2+ to Mn2+ and MnO2 respectively. Since MnO2 is a brown precipitate, it is impossible to observe the end-point if the titration is carried out in an alkaline medium. Furthermore, both Fe2+ and Fe3+ form coloured and colourless complexes with Cl– and PO43- respectively. As it is easier to observe the pale pink end-point in a colourless solution than a yellow solution, H3PO4 is preferred.

 

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Equivalence point vs End point

What is the difference between equivalence point and end point?

The equivalence point (also known as stoichiometric point) is a theoretical point during titration where the number of moles of an analyte is stoichiometrically equal to the number of moles of a solution of known concentration. For example, in the titration of 0.1 M HCl with 20 cm3 of 0.2 M NaOH, the equivalence point is when 40 cm3 of the acid is added to the base.

The end point, however, is the point during titration when a colour change of the solution in the flask is estimated by the person performing the titration. It may not coincide with the equivalence point. Even when instruments like colorimeters are used, the end point may not be exactly equal to the equivalence point due to a limit in the sensitivity of the instrument. However, the deviation between the end point and the equivalence point of a titration is usually very small and does not affect calculations significantly, whether the point is recorded through human observation or by a machine.

 

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pH indicators

A pH indicator is usually a weak organic acid that changes colour according to the pH of the solution to which it is added. As such, it serves as a useful marker of the pH of a solution, especially during the course of a titration.

Being a weak acid, a pH indicator HIn partially dissociates in a solution to give H+ and its conjugate base In–, with the undissociated acid and its conjugate base having different colours.

In a titration experiment, a few drops of a pH indicator, e.g. methyl orange (see below diagram), is added to the solution in the conical flask. If the solution is an acid, methyl orange, according to Le Chatelier’s principle, will be predominantly in the undissociated form, resulting in the solution being red.

As a base is gradually added to the solution, the above equilibrium shifts to the right. With more conjugate base formed, the mixture of HIn and In– makes the solution orange, which for a strong acid-strong base titration indicates that the end point is very near. Thereafter, a couple of drops of base turn the solution yellow and complete the titration.

The following table lists some common indicators that are suitable for various acid-base titrations:

Burette

 Conical

flask

 Indicator Colour change

Strong acid

Strong base

Methyl orange

Yellow → Red

Strong base

Strong acid

Red → Yellow

Strong acid

Weak base

Methyl red

Yellow → Red

Weak base

Strong acid

Red → Yellow

Strong base

Weak acid

Phenolphthalein

Colorless → Purple-pink
Weak acid Strong base

Purple-pink → Colorless

If you are interested to find out more about how to choose an appropriate pH indicator, you can read this article and its related articles in the intermediate section.

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